Question

The first ionization potential of \(\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}\) and \(\mathrm{Si}\) are in the order: (a) \(\mathrm{Na}>\mathrm{Mg}>\mathrm{Al}<\mathrm{Si}\) (b) \(\mathrm{Na}>\mathrm{Mg}>\mathrm{Al}>\mathrm{Si}\) (c) \(\mathrm{Na}<\mathrm{Mg}<\mathrm{AKSi}\) (d) \(\mathrm{Na}<\mathrm{Mg}>\mathrm{AKSi}\)

Short answer

Option (d) is correct: \\( \mathrm{Na}<\mathrm{Mg}>\mathrm{Al}<\mathrm{Si} \\).

Step-by-step solution

Understanding Ionization Potential

Ionization potential (IP) is the energy required to remove an electron from an atom in its gaseous state. Generally, elements on the right side of the periodic table have higher ionization potentials compared to those on the left. Also, in a period, the IP increases across from left to right due to increasing nuclear charge.

Analyze Periodic Trends

Sodium (Na), Magnesium (Mg), Aluminum (Al), and Silicon (Si) are all in the same period, i.e., Period 3. Generally, across a period from left to right, ionization potential increases.

Compare Specific Elements

Comparing Na, Mg, Al, and Si, Na typically has the smallest IP as it is the furthest left, making it easy to remove an electron. Mg has a higher IP than Na because it fulfills a filled s-orbital which provides some extra stability. Al has one additional electron in the p-orbital which is easier to remove, generally making its IP less than Mg. Si, having more protons and electrons, typically has a higher IP, potentially higher than both Na and Al.

Account for Exceptions

While the general trend is for increasing IP across a period, there can be exceptions due to electron configurations. Mg (with a full s-orbital) has a higher IP than both Al and Si because removing the p-orbital electron in Al is easier than removing from a full s-orbital in Mg.

Key concepts

Periodic Trends

In the periodic table, a fascinating pattern emerges known as periodic trends. These are specific patterns observed in the chemical properties of the elements. One key trend to note is that as you move from left to right across a period, the ionization potential generally increases. This is because the atomic number increases, leading to a greater positive charge in the nucleus. Consequently, the added electrons are more strongly attracted to the nucleus, requiring more energy to be removed.

There are, however, some exceptions due to unique electron configurations which may cause deviations from the expected trends. For example, elements with filled or half-filled sublevels tend to have slightly higher ionization energies, due to increased stability. Understanding these trends and exceptions can help you predict the behavior of different elements in chemical reactions.

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. Each element has a unique electron configuration that determines its chemical behavior. For instance, the electron configuration of Magnesium (Mg) is \[1s^22s^22p^63s^2\]. This full s-orbital provides some extra stability to Mg.

Naturally, elements like Sodium (Na) which sequentially fill these orbitals generally have lower ionization energy due to their position in the configuration. Recognizing the electron arrangement helps in understanding the chemical properties and potential reactions of each element. It explains why some electrons are easier to remove, impacting the ionization energy of the element. This mechanism is crucial for understanding concepts like periodic trends and predicting how elements will behave.

Periodic Table

The periodic table is a powerful tool that organizes all known elements in a structured way, highlighting their properties. Elements are arranged in order of increasing atomic number. The table is divided into periods and groups, allowing scientists and students to quickly understand the traits of each element.

Periods are the horizontal rows, and as you progress across them, you experience a change in periodic trends like ionization energy. Groups, the vertical columns, contain elements with similar chemical properties due to the same number of valence electrons. Utilizing the periodic table, one can identify how elements will interact, their stability, and reactivity patterns. It’s the roadmap for a chemist, unlocking the secrets of elemental relationships.

First Ionization Energy

First ionization energy is the amount of energy needed to remove the most loosely bound electron from a gaseous atom. This energy varies between elements and is structured by specific trends in the periodic table. Generally, ionization energy increases across a period and decreases down a group due to the increasing atomic radius.

For instance, Sodium (Na) has a lower ionization energy compared to Magnesium (Mg), Silicon (Si), and Aluminum (Al). This is due to its position in the period and having only one electron in its outermost shell, making it easier to remove. Meanwhile, Mg stands out due to its filled s-orbital, which grants it greater stability and consequently, a higher ionization energy than one might expect. Understanding this concept helps explain the reactivity and chemical behavior of elements.